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Chapter 3.2 Formal Charges

Chapter 3.2 Formal Charges

Postby josybabe2421 » Sat Aug 16, 2014 8:00 pm

I'm so confused. section 3.2 in the text book talks about Formal Charges and how they can be calculated by taking the number of valence shell electrons, minus the number of non bonding electrons, minus half of the bonding electrons. The examples in the book look fantastic but whenever I try it on another molecule, I can't get it to work.

For example, question 13 in the lesson practice test 1 for this chapter. The explanation for that question just gives another way of calculating formal charge which isn't in the text book.

I know I'm making some major silly mistake, but I have no idea what it is. Could someone explain to me how to calculate the formal charge on Sulfur in question 13 using the method mentioned on page CHM-27 of the GAMSAT textbook? Its driving me nuts.

Thank you so much :)
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Re: Chapter 3.2 Formal Charges

Postby goldstanda3269 » Mon Aug 18, 2014 5:33 pm

"section 3.2 in the text book talks about Formal Charges and how they can be calculated by taking the number of valence shell electrons, minus the number of non bonding electrons, minus half of the bonding electrons."
Applying the VSEPR rules for sulfur (CHM 3.2, 3.3):
the number of valence shell electrons: 6 (just like oxygen since it is right below oxygen in the periodic table)
minus the number of non bonding electrons: - 2 (see diagram)
minus half of the bonding electrons: - (1/2) 6
= -1


There are different ways to calculate formal charge so you should try to stay flexible. Here is the good news: you can almost always get away with using the simplest method as described in the Explanation. If they want you to use a more complex method (VSEPR, as above) then they will remind you of all the rules just to see if you can apply those rules to a problem.
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Re: Chapter 3.2 Formal Charges

Postby josybabe2421 » Mon Aug 18, 2014 6:09 pm

Thank you for your help :) I think I understand where I went wrong now.

I didn't think that Sulfur could share 2 of its own electrons to form a bond with Oxygen without Oxygen sharing any of its electrons. For some reason I had it in my head each atom needed to share one of its electrons to form a bond.

So I thought there must actually be 3 electrons left and began to think I was reading the diagram wrong and calculated it as being 0.

This seems to be a mistake I've made pretty consistently, chemistry is definitely my weakest subject, I will have to study it twice as hard :) thank you so much for your help.
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Re: Chapter 3.2 Formal Charges

Postby goldstanda3269 » Tue Aug 19, 2014 7:10 am

just FYI: a regular covalent bond has contributions from both atoms; a "coordinate covalent" bond has electrons from only one atom (an example was given in the General Chemistry videos).
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Re: Chapter 3.2 Formal Charges

Postby josybabe2421 » Tue Aug 19, 2014 8:25 pm

Thank you :D I must have missed that. I will look it up in the video :)
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Re: Chapter 3.2 Formal Charges

Postby Lavi » Wed May 11, 2016 7:39 am

I'm a bit confused about working out the number of lone pairs from the structure provided in Q5 from the chapter review question-

5) For the incomplete Lewis structure drawn, how many lone pairs of electrons are present in the following molecule?

http://www.oatbooks.com/images/index_c72.png

A) 0
B) 1
C) 2
D) 3
E) 4

I keep seeing 3 rather than 4
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Re: Chapter 3.2 Formal Charges

Postby goldstanda3269 » Thu May 26, 2016 6:35 am

Lavi wrote:I'm a bit confused about working out the number of lone pairs from the structure provided in Q5 from the chapter review question-

5) For the incomplete Lewis structure drawn, how many lone pairs of electrons are present in the following molecule?
.....

I keep seeing 3 rather than 4


I think you will find either of these 2 images helpful:

Same as the practice question: http://session.masteringchemistry.com/p ... 07.079.jpg

Resonance forms: http://chemiris.chem.binghamton.edu/che ... sofig1.gif
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